Determination of R: The Gas Law Constant
Introduction:The main objective of this experiment is to better understand how well real gases obey
the ideal gas law and how to determine R: The idea gas low constant. At room temperature and
atmospheric pressure most gasses will obey the idea gas equation, PV=nRT, butthere aresmall
deviations from this law due to real gas molecules being finite in size and having mutual attractive
forces. These two causes for deviation are taken into account by the van der Waals equation
(P+((n2a)/V2))(V-nb)=nRT. This equation is applicable over a wider range of temperatures and
pressures than the ideal gas law. We will determine the value of R in its most common form Latm/mol-K by using both equations and from measured values of pressure, temperature, volume,
and moles of an enclosed sample of oxygen (prepared by the decomposition of potassium chlorate,
using manganese (IV) oxide as a catalyst) then we will do an error analysis on the experimentally
determined constant. 2KClO3(s)—MnO2(s)—>2KCl(s)+3O2(g)
P.S KClO3 is very dangerous, do not apply to much friction when scooping out powder, and do not
allow any of the chemicals to come into contact with the rubber stopper or a severe explosion may
occur.
Materials:
Balance
Bunsen Burner and hose
Test Tube
250 mL beaker
8 oz wide mouth bottle
Rubber stoppers
Pinch clamp
Clamp
Barometer (or readings)
Glass tubing
Rubber tubing
Thermometer
Ring Stand
Procedure: 1.Pre-weigh a dry beaker. Addabout 0.02g of Mno2 and about 0.3g of KClO3 to a test
tube and weigh to the nearest 0.001g.
2.Add about 200ml of distilled water into the bottle.
3.Assemble the apparatus but do not attach test tube, and be sure that tube B does not extend
below the water level in the bottle
4.Attach rubber bulb to tube B and apply pressure through it so that water will travel into tube A
then close the clamp when Tube A is filled.
5.Mix the solids in the test tube by rotating it, making sure none of the mixture is being lost from the
tube, and attach tube B.
6.Fill beaker about half full of water, insert tube A in it and open the pinch clamp. Lift the beaker
until the water levels are identical (to produce atmospheric pressure inside the bottle and test tube).
Close clamp and discard water in the beaker, dry beaker.
7.Set beaker with tube A on bench and open pinch clamp. A little water should flow, but should stop
quickly and tube A will remain filled. If that is not the case, check for leaks and start over.
8.Heat bottom of test tube gently so that a slow but steady stream of gas is produced, shown by the
flow of water into the beaker. When the evolution of gas slows a lot, increase the rate of heating,
then then until no more oxygen is produced. Let the apparatus cool and then equalize the water
levels in the beaker and the bottle as before and close the clamp.
9.Weigh the beaker with the water and get a temperature of the water. Using the temperature,
match the density and calculate the volume of water that was displaced with the chart. This is the
volume of oxygen produced. Round to temperature to the nearest whole. Record the barometric
pressure, the vapor pressure of the water can be found by the temperature of the waterfrom the
chart.
10.Remove test tube and weigh the tube and contents. The difference in mass from the original
mass of tube and contents is the mass of oxygen produced.
11.Return the tube to the instructor for proper disposal...
Results: On lab papers
Analysis of Results: The value for R that we obtained was very close to the accepted value of R, and
the values we obtained for R from both equations were nearly equal. From this we can presume that
both the van der Waals and the ideal gas law are pretty accurate for their purposes. Some sources of
error, and cause for difference from our R and the R constant, can be found when considering
slightly off measurements and rounding. The scale or table rounding could be off a bit. Some oxygen
might have dissolved in the water, causing the volume measurements to be off. Lastly, the constant
R might have been the
average constant found from the results of this experiment done numerous amounts of times.
Conclusion:We accept that most gasses, at least oxygen, will behave very closely to the ideal gas law
and the van der Waals equation, and that these two equations will give nearly the same results, with
the van der Waals a bit more accurate. We also accept that the ideal gas law constant R is correct.